A guide for the preparation and use of buffers in biological systems

Buffers
A guide for the preparation and use of
buffers in biological systems
Advancing your life science discoveries

Buffers
A guide for the preparation and use of
buffers in biological systems
By
Chandra Mohan, Ph.D.
A brand of EMD Biosciences, Inc.
Copyright © 2003 EMD Biosciences, Inc., An Affiliate of Merck KGaA, Darmstadt, Germany.
All Rights Reserved.

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ii

Table of Contents:
Why Does Calbiochem® Biochemicals Publish a Booklet on Buffers? . . . . . . . . . .1
Water, The Fluid of Life . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .2
Ionization of Water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .3
Dissociation Constants of Weak Acids and Bases . . . . . . . . . . . . . . . . . . . . . . . . . .4
Henderson-Hasselbach Equation: pH and pK a
. . . . . . . . . . . . . . . . . . . . . . . . . . . . .5
Determination of pK a
. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .6
pK a
Values for Commonly Used Biological Buffers . . . . . . . . . . . . . . . . . . . . . . . . .7
Buffers, Buffer Capacity, and Range . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .8
Biological Buffers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .10
Buffering in Cells and Tissues . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .10
Effect of Temperature on pH . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .12
Effect of Buffers on Factors Other than pH . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .13
Use of Water-Miscible Organic Solvents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .14
Solubility Equilibrium: Effect of pH on Solubility . . . . . . . . . . . . . . . . . . . . . . . .14
pH Measurements: Some Useful Tips . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .15
Choosing a Buffer . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .16
Preparation of Some Common Buffers for Use
in Biological Systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .18
Commonly Used Buffer Media in Biological Research . . . . . . . . . . . . . . . . . . . . .22
Isoelectric Point of Selected Proteins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .24
Isoelectric Point of Selected Plasma Proteins . . . . . . . . . . . . . . . . . . . . . . . . . . . .26
Approximate pH and Bicarbonate Concentration in
Extracellular Fluids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .26
Ionization Constants K and pK a
for Selected Acids and
Bases in Water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .27
Physical Properties of Some Commonly Used Acids . . . . . . . . . . . . . . . . . . . . . . .27
Some Useful Tips for Calculation of Concentrations and
Spectrophotometric Measurements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .28
CALBIOCHEM® Buffers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .30
iii

Why Does Calbiochem® Biochemicals Publish a
Booklet on Buffers?
We are frequently asked questions on the use of buffers that we offer to research
laboratories. This booklet is designed to help answer several basic questions
about the use of buffers in biological systems. The discussion presented here is
by no means complete, but we hope it will help in the understanding of general
principles involved in the use of buffers.
Almost all biological processes are pH dependent. Even a slight change in pH can
result in metabolic acidosis or alkalosis, resulting in severe metabolic complications.
The purpose of a buffer in biological system is to maintain intracellular
and extracellular pH within a very narrow range and resist changes in pH in the
presence of internal and external influences. Before we begin a discussion of
buffers and how they control hydrogen ion concentrations, a brief explanation
of the role of water and equilibrium constants of weak acids and bases is
necessary.
1

Water: The Fluid of Life
Water constitutes about 70% of the mass of most living creatures. All biological
reactions occur in an aqueous medium. All aspects of cell structure and function
are adapted to the physical and chemical properties of water. Hence, it is
essential to understand some basic properties of water and its ionization
products, i.e., H + and OH¯. Both H + and OH¯ influence the structure, assembly,
and properties of all macromolecules in the cell.
Water is a polar solvent that dissolves most charged molecules. Water dissolves
most salts by hydrating and stabilizing the cations and anions by weakening
their electrostatic interactions (Figure 1). Compounds that readily dissolve in
water are known as HYDROPHILIC compounds. Nonpolar compounds such as
chloroform and ether do not interact with water in any favorable manner and are
known as HYDROPHOBIC compounds. These compounds interfere with
hydrogen bonding among water molecules.
Figure 1: Electrostatic interaction of Na + and Cl¯ ions and water molecules.
Several biological molecules, such as protein, certain vitamins, steroids, and
phospholipids contain both polar and nonpolar regions. They are known as
AMPHIPATHIC molecules. The hydrophilic region of these molecules are
arranged in a manner that permits maximum interaction with water molecules.
However, the hydrophobic regions assemble together exposing only the smallest
area to water.
2

Ionization of Water
Water molecules undergo reversible ionization to yield H + and OH¯ as per the
following equation.
H 2
O → ← H + + OH¯
The degree of ionization of water at equilibrium is fairly small and is given by
the following equation where K eq
is the equilibrium constant.
K eq
[H + ][OH¯]
= ______________
[H 2
O]
At 25°C, the concentration of pure water is 55.5 M (1000 ÷ 18; M.W. 18.0).
Hence, we can rewrite the above equation as follows:
K eq
[H + ][OH¯]
= ______________
55.5 M
or
(55.5)(K eq
) = [H+][OH¯]
For pure water electrical conductivity experiments give a K eq
value of 1.8 x
10 -16 M at 25°C.
Hence,
(55.5 M)(1.8 x 10 -16 M) = [H + ][OH¯]
or
99.9 x 10 -16 M 2 = [H + ][OH¯]
or
1.0 x 10 -14 M 2 = [H + ][OH¯]
[H + ][OH¯], ion product of water, is always equal to 1.0 x 10 -14 M 2 at 25°C. When
[H + ] and [OH¯] are present in equal amounts then the solution gives a neutral pH.
Here [H + ][OH¯] = [H + ] 2
or
[H + ] = 1 x 10 -14 M 2
and
[H + ] = [OH¯] = 10 -7 M
As the total concentration of H + and OH¯ is constant, an increase in one ion is
compensated by a decrease in the concentration of other ion. This forms the
basis for the pH scale.
3

Dissociation Constants of Weak Acids and Bases
Strong acids (hydrochloric acid, sulfuric acid, etc.) and bases (sodium hydroxide,
potassium hydroxide, etc.) are those that are completely ionized in dilute
aqueous solutions.
In biological systems one generally encounters only weak acids and bases. Weak
acids and bases do not completely dissociate in solution. They exist instead as an
equilibrium mixture of undissociated and dissociated species. For example, in
aqueous solution, acetic acid is an equilibrium mixture of acetate ion, hydrogen
ion, and undissociated acetic acid. The equilibrium between these species can be
expressed as:
CH 3
COOH
k 1
→ ←
H + + CH 3
COO¯
k 2
where k 1
represents the rate constant of dissociation of acetic acid to acetate and
hydrogen ions, and k 2
represents the rate constant for the association of acetate
and hydrogen ions to form acetic acid. The rate of dissociation of acetic acid,
-d[CH 3
COOH ]/dt, is dependent on the rate constant of dissociation (k 1
) and the
concentration of acetic acid [CH 3
COOH] and can be expressed as:
____________________
d [CH 3
COOH]
dt
= k 1
[CH 3
COOH]
Similarly, the rate of association to form acetic acid, d[HAc]/dt, is dependent on
the rate constant of association (k 2
) and the concentration of acetate and
hydrogen ions and can be expressed as:
d [CH 3
COOH ]
__________________
dt
= k 2
[H + ] [CH 3
COO¯]
Since the rates of dissociation and reassociation are equal under equilibrium
conditions:
k 1
[CH 3
COOH ] = k 2
[H + ] [CH 3
COO¯]
or
k 1
[H + ] [CH 3
COO¯]
_______ = ____________________
[CH 3
COOH]
and
where
4
k 2
k _______ 1
k 2
[H + ] [CH 3
COO¯]
K a
= ___________________
[CH 3
COOH]
=K a
(Equilibrium constant)

This equilibrium expression can now be rearranged to
[CH 3
COOH]
[H + ] = K a
_______________
[CH 3
COO¯]
where the hydrogen ion concentration is expressed in terms of the equilibrium
constant and the concentrations of undissociated acetic acid and acetate ion. The
equilibrium constant for ionization reactions is called the ionization constant or
dissociation constant.
Henderson-Hasselbach Equation: pH and pK a
The relationship between pH, pK a
, and the buffering action of any weak acid and
its conjugate base is best explained by the Henderson-Hasselbach equation. In
biological experiments, [H + ] varies from 10 -1 M to about 10 -10 M. S.P.L.
Sorenson, a Danish chemist, coined the “p” value of any quantity as the negative
logarithm of the hydrogen ion concentration. Hence, for [H + ] one can write the
following equation:
pH = – log [H + ]
Similarly pK a
can be defined as – log K a
. If the equilibrium expression is
converted to – log then
[CH 3
COOH]
– log [H + ] = – log K a
– log ______________
[CH 3
COO¯]
and pH and pK a
substituted:
[CH 3
COOH]
pH = pK a
– log ________________
[CH 3
COO¯]
pH =
or
[CH 3
COO¯]
pK a
+ log _______________
[CH 3
COOH]
When the concentration of acetate ions equals the concentration of acetic acid,
log [CH 3
COO¯]/[CH 3
COOH] approaches zero (the log of 1) and pH equals pK a
(the
pK a
of acetic acid is 4.745). Acetic acid and acetate ion form an effective
buffering system centered around pH 4.75. Generally, the pK a
of a weak acid or
base indicates the pH of the center of the buffering region.
The terms pK and pK a
are frequently used interchangeably in the literature. The
term pK a
(“a” refers to acid) is used in circumstances where the system is being
considered as an acid and in which hydrogen ion concentration or pH is of
5

interest. Sometimes the term pK b
is used. pK b
(“b” refers to base) is used when the
system is being considered as a base and the hydroxide ion concentration or pOH
is of greater interest.
Determination of pK a
pKa values are generally determined by titration. A carefully calibrated,
automated, recording titrator is used, the free acid of the material to be measured
is titrated with a suitable base, and the titration curve is recorded. The pH of the
solution is monitored as increasing quantities of base are added to the solution.
Figure 2 shows the titration curve for acetic acid. The point of inflection
indicates the pK a
value. Frequently, automatic titrators record the first derivative
of the titration curve, giving more accurate pK a
values.
Polybasic buffer systems can have more than one useful pK a
value. Figure 3
shows the titration curve for phosphoric acid, a tribasic acid. Note that the curve
has five points of inflection. Three indicate pK a1
, pK a2
and pK a3
, and two
additional points indicate where H 2
PO 4 – and HPO 4 – exist as the sole species.
8
6
pK a = 4.76
pH
4
2
0
NaOH
Figure 2: Titration Curve for Acetic Acid
12
pK a3 = 12.32
10
8
pH
6
pK a2 = 7.21
4
2
pK a1 = 2.12
NaOH
Figure 3: Titration Curve for Phosphoric Acid
6

Table 1: pK a
Values for Commonly Used Biological Buffers and Buffer Constituents
Product Cat. No. M.W.
ADA, Sodium Salt 114801 212.2 6.60
2-Amino-2-methyl-1,3-propanediol 164548 105.1 8.83
BES, ULTROL® Grade 391334 213.2 7.15
Bicine, ULTROL® Grade 391336 163.2 8.35
BIS-Tris, ULTROL® Grade 391335 209.2 6.50
BIS-Tris Propane, ULTROL® Grade 394111 282.4 6.80
Boric Acid, Molecular Biology Grade 203667 61.8 9.24
Cacodylic Acid 205541 214.0 6.27
CAPS, ULTROL® Grade 239782 221.3 10.40
CHES, ULTROL® Grade 239779 207.3 9.50
Citric Acid, Monohydrate, Molecular Biology Grade 231211 210.1 4.76
Glycine 3570 75.1 2.34 1
Glycine, Molecular Biology Grade 357002 75.1 2.34 1
Glycylglycine, Free Base 3630 132.1 8.40
HEPES, Free Acid, Molecular Biology Grade 391340 238.3 7.55
HEPES, Free Acid, ULTROL® Grade 391338 238.3 7.55
HEPES, Free Acid Solution 375368 238.3 7.55
HEPES, Sodium Salt, ULTROL® Grade 391333 260.3 7.55
HEPPS, ULTROL® Grade 391339 252.3 8.00
Imidazole, ULTROL® Grade 4015 68.1 7.00
MES, Free Acid, ULTROL® Grade 475893 195.2 6.15
MES, Sodium Salt, ULTROL® Grade 475894 217.2 6.15
MOPS, Free Acid, ULTROL® Grade 475898 209.3 7.20
MOPS, Sodium Salt, ULTROL® Grade 475899 231.2 7.20
PIPES, Free Acid, Molecular Biology Grade 528133 302.4 6.80
PIPES, Free Acid, ULTROL® Grade 528131 302.4 6.80
PIPES, Sodium Salt, ULTROL® Grade 528132 325.3 6.80
PIPPS 528315 330.4 3.73 2
Potassium Phosphate, Dibasic, Trihydrate, Molecular Biology Grade 529567 228.2 7.21 3
Potassium Phosphate, Monobasic 529565 136.1 7.21 3
Potassium Phosphate, Monobasic, Molecular Biology Grade 529568 136.1 7.21 3
Sodium Phosphate, Dibasic 567550 142.0 7.21 3
Sodium Phosphate, Dibasic, Molecular Biology Grade 567547 142.0 7.21 3
Sodium Phosphate, Monobasic 567545 120.0 7.21 3
Sodium Phosphate, Monobasic, Monohydrate, Molecular Biology Grade 567549 138.0 7.21 3
TAPS, ULTROL® Grade 394675 243.2 8.40
TES, Free Acid, ULTROL® Grade 39465 229.3 7.50
TES, Sodium Salt, ULTROL® Grade 394651 251.2 7.50
Tricine, ULTROL® Grade 39468 179.2 8.15
Triethanolamine, HCl 641752 185.7 7.66
Tris Base, Molecular Biology Grade 648310 121.1 8.30
Tris Base, ULTROL® Grade 648311 121.1 8.30
Tris, HCl, Molecular Biology Grade 648317 157.6 8.30
Tris, HCl, ULTROL® Grade 648313 157.6 8.30
Trisodium Citrate, Dihydrate 567444 294.1 —
Trisodium Citrate, Dihydrate, Molecular Biology Grade 567446 294.1 —
1. pK a1
= 2.34; pK a2
= 9.60
2. pK a1
= 3.73; pK a2
= 7.96 (100 mM aqueous solution, 25°C).
3. Phosphate buffers are normally prepared from a combination of the monobasic and dibasic salts, titrated against
each other to the correct pH. Phosphoric acid has three pK a
values: pK a1
= 2.12; pK a2
= 7.21; pK a3
= 12.32
pK a
at 20°C
7

Buffers, Buffer Capacity and Range
Buffers are aqueous systems that resist changes in pH when small amounts of
acid or base are added. Buffer solutions are composed of a weak acid (the proton
donor) and its conjugate base (the proton acceptor). Buffering results from two
reversible reaction equilibria in a solution wherein the concentration of proton
donor and its conjugate proton acceptor are equal. For example, in a buffer
system when the concentration of acetic acid and acetate ions are equal, addition
of small amounts of acid or base do not have any detectable influence on the pH.
This point is commonly known as the isoelectric point. At this point there is no
net charge and pH at this point is equal to pK a
.
[CH3COO¯]
pH = pK a
+ log ________________
[CH3COOH]
At isoelectric point [CH 3
COO¯] = [CH 3
COOH] hence, pH = pK a
Buffer capacity is a term used to describe the ability of a given buffer to resist
changes in pH on addition of acid or base. A buffer capacity of 1 is when 1 mol
of acid or alkali is added to 1 liter of buffer and pH changes by 1 unit. The buffer
capacity of a mixed weak acid-base buffer is much greater when the individual
pK a
values are in close proximity with each other. It is important to note that the
buffer capacity of a mixture of buffers is additive.
Buffers have both intensive and extensive properties. The intensive property is a
function of the pK a
value of the buffer acid or base. Most simple buffers work
effectively in the pH scale of pK a
± 1.0. The extensive property of the buffers is
also known as the buffer capacity. It is a measure of the protection a buffer offers
against changes in pH. Buffer capacity generally depends on the concentration
of buffer solution. Buffers with higher concentrations offer higher buffering
capacity. On the other hand, pH is dependent not on the absolute concentrations
of buffer components but on their ratio.
Using the above equation we know that when pH = pK a
the concentrations of
acetic acid and acetate ion are equal. Using a hypothetical buffer system of HA
(pK a
= 7.0) and [A – ], we can demonstrate how the hydrogen ion concentration,
[H + ], is relatively insensitive to external influence because of the buffering
action.
For example:
If 100 ml of 10 mM (1x 10 -2 M) HCl are added to 1.0 liter of 1.0 M NaCl at pH 7.0,
the hydrogen ion concentration, [H + ], of the resulting 1.1 liter of solution can be
calculated by using the following equation:
8

[H + ] x Vol = [H + ] o
x Vol o
where
Vol o
= initial volume of HCl solution (in liters)
[H + ] o
= initial hydrogen ion concentration (M)
Vol = final volume of HCl + NaCl solutions (in liters)
[H + ] = final hydrogen ion concentration of HCl + NaCl solution (M)
Solving for [H + ]:
[H + ] x 1.1 liter = 1.0 x 10 -2 x 0.1 = 1 x 10 -3
[H + ] = 9.09 x 10 -4
or pH = 3.04
Thus, the addition of 1.0 x 10 -3 mol of hydrogen ion resulted in a pH change of
approximately 4 pH units (from 7.0 to 3.04).
If a buffer is used instead of sodium chloride, a 1.0 M solution of HA at pH 7.0
will initially have:
[HA] = [A] = 0.5 M
[A]
pH = pK + log ______
[HA]
0.5
pH = 7.0 + log ______ 0.5
or pH = 7.0
When 100 ml of 1.0 x 10 -2 M (10 mM) HCl is added to this system, 1.0 x 10 -3 mol
of A – is converted to 1.0 x 10 -3 mol of HA, with the following result:
0.499/1.1
pH = 7.0 + log _______________
0.501/1.1
pH = 7.0 - 0.002 or pH = 6.998
Hence, it is clear that in the absence of a suitable buffer system there was a pH
change of 4 pH units, whereas in a buffer system only a trivial change in pH was
observed indicating that the buffer system had successfully resisted a change in
pH. Generally, in the range from [A]/[HA] = 0.1 to [A]/[HA] = 10.0, effective
buffering exists. However, beyond this range, the buffering capacity may be
significantly reduced.
9

Biological Buffers
Biological buffers should meet the following general criteria:
• Their pK a
should reside between 6.0 to 8.0.
• They should exhibit high water solubility and minimal solubility in organic
solvents.
• They should not permeate cell membranes.
• They should not exhibit any toxicity towards cells.
• The salt effect should be minimum, however, salts can be added as required.
• Ionic composition of the medium and temperature should have minimal
effect of buffering capacity.
• Buffers should be stable and resistant to enzymatic degradation.
• Buffer should not absorb either in the visible or in the UV region.
Most of the buffers used in cell cultures, isolation of cells, enzyme assays, and
other biological applications must possess these distinctive characteristics.
Good's zwitterionic buffers meet these criteria. They exhibit pK a
values at or near
physiological pH. They exhibit low interference with biological processes due to
the fact that their anionic and cationic sites are present as non-interacting
carboxylate or sulfonate and cationic ammonium groups respectively.
Buffering in Cells and Tissues
A brief discussion of hydrogen ion regulation in biological systems highlights
the importance of buffering systems. Amino acids present in proteins in cells and
tissues contain functional groups that act as weak acid and bases. Nucleotides
and several other low molecular weight metabolites that undergo ionization also
contribute effectively to buffering in the cell. However, phosphate and bicarbonate
buffer systems are most predominant in biological systems.
The phosphate buffer system has a pK a
of 6.86. Hence, it provides effective
buffering in the pH range of 6.4 to 7.4. The bicarbonate buffer system plays an
important role in buffering the blood system where in carbonic acid acts as a
weak acid (proton donor) and bicarbonate acts as the conjugate base (proton
acceptor). Their relationship can be expressed as follows:
[H + ][HCO 3¯]
K 1
= ______________
[H 2
CO 3
]
In this system carbonic acid (H 2
CO 3
) is formed from dissolved carbon dioxide
and water in a reversible manner. The pH of the bicarbonate system is dependent
on the concentration of carbonic acid and bicarbonate ion. Since carbonic acid
10

concentration is dependent upon the amount of dissolved carbon dioxide the
ultimate buffering capacity is dependent upon the amount of bicarbonate and
the partial pressure of carbon dioxide.
+ _
H + HCO 3
H 2
CO 2
H 2
OCO 2
H 2
O
CO 2
Blood
Lung
Air Space
Figure 4: Relationship between bicarbonate buffer system and carbon dioxide.
In air breathing animals, the bicarbonate buffer system maintains pH near 7.4.
This is possible due to the fact that carbonic acid in the blood is in equilibrium
with the carbon dioxide present in the air. Figure 4 highlights the mechanism
involved in blood pH regulation by the bicarbonate buffer system. Any increase
in partial pressure of carbon dioxide (as in case of impaired ventilation) lowers
the ratio of bicarbonate to pCO 2
resulting in a decrease in pH (acidosis). The
acidosis is reversed gradually when kidneys increase the absorption of bicarbonate
at the expense of chloride. Metabolic acidosis resulting from the loss of
bicarbonate ions (such as in severe diarrhea or due to increased keto acid
formation) leads to severe metabolic complications warranting intravenous
bicarbonate therapy.
During hyperventilation, when excessive amounts of carbon dioxide are
eliminated from the system (thereby lowering the pCO 2
), pH of the blood
increases resulting in alkalosis. This is commonly seen in conditions such as
pulmonary embolism and hepatic failure. Metabolic alkalosis generally results
when bicarbonate levels are higher in the blood. This is commonly observed after
vomiting of acidic gastric secretions. Kidneys compensate for alkalosis by
increasing the excretion of bicarbonate ions. However, an obligatory loss of
sodium occurs under these circumstances.
In case of severe alkalosis the body is depleted of water, H + , Cl¯ and to some
extent Na + . A detailed account of metabolic acidosis and alkalosis is beyond the
scope of this booklet. Readers are advised to consult a suitable text book of
physiology for more detailed information on the mechanisms involved.
11

Effect of Temperature on pH
Generally when we consider the use of buffers we make following two assumptions.
(a) The activity coefficients of the buffer ions is approximately equal to 1
over the useful range of buffer concentrations
(b) The value of K a
is constant over the working range of temperature.
However, in real practice one observes that pH changes slightly with change in
temperature. This might be very critical in biological systems where a precise
hydrogen ion concentration is required for reaction systems to operate with
maximum efficiency. Figure 5 presents the effect of temperature on the pH of
phosphate buffer. The difference might appear to be slight but it has significant
biological importance. Although the mathematical relationship of activity and
temperature may be complicated, the actual change of pK a
with temperature
(∆pK a
/°C) is approximately linear. Table 2 presents the pK a
and ∆pK a
/°C for several
selected zwitterionic buffers commonly used in biological experimentation.
6.7
6.8
pH
6.9
7.0
0 10 20 30 40
Temperature, ºC
Figure 5: Effect of Temperature on pH of Phosphate Buffer
12

Table 2: pK a
and DpK a
/°C of Selected Buffers
Buffer
M.W.
pK a
pK a
DpK a
/°C Binding to
(20°C) (37°C) Metal Ions
MES 195.2 6.15 5.97 -0.011 Negligible metal ion binding
ADA 212.2 6.60 6.43 -0.011 Cu 2+ , Ca 2+ , Mn 2+ . Weaker
binding with Mg 2+ .
BIS-Tris Propane* 282.4 6.80 — -0.016 —
PIPES 302.4 6.80 6.66 -0.009 Negligible metal ion binding
ACES 182.2 6.90 6.56 -0.020 Cu 2+ . Does not bind Mg 2+ ,
Ca 2+ , or Mn 2+ .
BES 213.3 7.15 6.88 -0.016 Cu 2+ . Does not bind
Mg 2+ , Ca 2+ , or Mn 2+ .
MOPS 209.3 7.20 6.98 -0.006 Negligible metal ion binding
TES 229.3 7.50 7.16 -0.020 Slightly to Cu 2+ . Does not bind
Mg 2+ , Ca 2+ , or Mn 2+ .
HEPES 238.3 7.55 7.30 -0.014 None
HEPPS 252.3 8.00 7.80 -0.007 None
Tricine 179.2 8.15 7.79 -0.021 Cu 2+ . Weaker binding
with Ca 2+ , Mg 2+ , and Mn 2+ .
Tris* 121.1 8.30 7.82 -0.031 Negligible metal ion binding
Bicine 163.2 8.35 8.04 -0.018 Cu 2+ . Weaker binding
with Ca 2+ , Mg 2+ , and Mn 2+ .
Glycylglycine 132.1 8.40 7.95 -0.028 Cu 2+ . Weaker
binding with Mn 2+ .
CHES 207.3 9.50 9.36 -0.009 —
CAPS 221.32 10.40 10.08 -0.009 —
* Not a zwitterionic buffer
Effects of Buffers on Factors Other than pH
It is of utmost importance that researchers establish the criteria and determine
the suitability of a particular buffer system. Some weak acids and bases may
interfere with the reaction system. For example, citrate and phosphate buffers are
not recommended for systems that are highly calcium-dependent. Citric acid and
its salts are powerful calcium chelators. Phosphates react with calcium producing
insoluble calcium phosphate that precipitates out of the system. Phosphate ions
in buffers can inhibit the activity of some enzymes, such as carboxypeptidase,
fumarease, carboxylase, and phosphoglucomutase.
Tris(hydroxy-methyl)aminomethane can chelate copper and also acts as a
competitive inhibitor of some enzymes. Other buffers such as ACES, BES, and
TES, have a tendency to bind copper. Tris-based buffers are not recommended
when studying the metabolic effects of insulin. Buffers such as HEPES and
HEPPS are not suitable when a protein assay is performed by using Folin
reagent. Buffers with primary amine groups, such as Tris, may interfere with the
Bradford dye-binding method of protein assay. Borate buffers are not suitable for
gel electrophoresis of protein, they can cause spreading of the zones if polyols
are present in the medium.
13

Use of Water-Miscible Organic Solvents
Most pH measurements in biological systems are performed in the aqueous
phase. However, sometimes mixed aqueous-water-miscible solvents, such as
methanol or ethanol, are used for dissolving compounds of biological importance.
These organic solvents have dissociation constants that are very low
compared to that of pure water or of aqueous buffers (for example, the dissociation
constant of methanol at 25°C is 1.45 x 10 -17 , compared to 1.0 x 10 -14 for
water). Small amounts of methanol or ethanol added to the aqueous medium will
not affect the pH of the buffer. However, even small traces of water in methanol
or DMSO can significantly change the pH of these organic solvents.
Solubility Equilibrium: Effect of pH on Solubility
A brief discussion of the effect of pH on solubility is of significant importance
when dissolution of compounds into solvents is under consideration. Changes in
pH can affect the solubility of partially soluble ionic compounds.
Example:
Here
Mg(OH) 2
→ ← Mg 2+ + 2OH¯
[Mg 2+ ] [OH¯ ] 2
K = ________________
[Mg(OH) 2
]
As a result of the common ion effect, the solubility of Mg(OH) 2
can be increased
or decreased. When a base is added the concentration of OH¯ increases and shifts
the solubility equilibrium to the left causing a diminution in the solubility of
Mg(OH) 2
. When an acid is added to the solution, it neutralizes the OH¯ and shifts
the solubility equilibrium to the right. This results in increased dissolution of
Mg(OH) 2
.
14

pH Measurements: Some Useful Tips
1. A pH meter may require a warm up time of several minutes. When a pH
meter is routinely used in the laboratory, it is better to leave it “ON” with the
function switch at “standby.”
2. Set the temperature control knob to the temperature of your buffer solution.
Always warm or cool your buffer to the desired temperature before checking
final pH.
3. Before you begin make sure the electrode is well rinsed with deionized water
and wiped off with a clean absorbent paper.
4. Always rinse and wipe the electrode when switching from one solution to
another.
5. Calibrate your pH meter by using at least two standard buffer solutions.
6. Do not allow the electrode to touch the sides or bottom of your container.
When using a magnetic bar to stir the solution make sure the electrode tip is
high enough to prevent any damage.
7. Do not stir the solution while taking the reading.
8. Inspect your electrode periodically. The liquid level should be maintained as
per the specification provided with the instrument .
9. Glass electrodes should not be left immersed in solution any longer than
necessary. This is important especially when using a solution containing
proteins. After several pH measurements of solutions containing proteins,
rinse the electrode in a mild alkali solution and then wash several times with
deionized water.
10. Water used for preparation of buffers should be of the highest possible purity.
Water obtained by a method combining deionzation and distillation is highly
recommended.
11. To avoid any contamination do not store water for longer than necessary. Store
water in tightly sealed containers to minimize the amount of dissolved gases.
12. One may sterile-filter the buffer solution to prevent any bacterial or fungal
growth. This is important when large quantities of buffers are prepared and
stored over a long period of time.
15

CHOOSING A BUFFER
1. Recognize the importance of the pK a
. Select a buffer that has a pK a
value
close to the middle of the range required. If you expect the pH to drop during
the experiment, choose a buffer with a pK a
slightly lower than the working
pH. This will permit the buffering action to become more resistant to changes
in hydrogen ion concentration as hydrogen ions are liberated. Conversely, if
you expect the pH to rise during the experiment, choose a buffer with a pK a
slightly higher than the working pH. For best results, the pK a
of the buffer
should not be affected significantly by buffer concentration, temperature,
and the ionic constitution of the medium.
2. Adjust pH at desired temperature. The pK a
of a buffer, and hence the pH,
changes slightly with temperature. It is best to adjust the final pH at the
desired temperature.
3. Prepare buffers at working conditions. Always try to prepare your buffer
solution at the temperature and concentration you plan to use during the
experiment. If you prepare stock solutions make dilutions just prior to use.
4. Purity and cost. Compounds used should be stable and be available in high
purity and at moderate cost.
5. Spectral properties: Buffer materials should have no significant absorbance
between 240 to 700 nm range.
6. Some weak acids (or bases) are unsuitable for use as buffers in certain
cases. Citrate and phosphate buffers are not suitable for systems that are
highly calcium-dependent. Citric acid and its salts are chelators of calcium
and calcium phosphates are insoluble and will precipitate out. Use of these
buffers may lower the calcium levels required for optimum reaction. Tris
(hydroxymethyl) aminomethane is known to chelate calcium and other
essential metals.
7. Buffer materials and their salts can be used together for convenient
buffer preparation. Many buffer materials are supplied both as a free acid
(or base) and its corresponding salt. This is convenient when making a series
of buffers with different pH’s. For example, solutions of 0.1 M HEPES and 0.1
M HEPES, sodium salt, can be mixed in an infinite number of ratios between
10:1 and 1:0 to provide 0.1 M HEPES buffer with pH values ranging from
6.55 to 8.55.
16

8. Use stock solutions to prepare phosphate buffers. Mixing precalculated
amounts of monobasic and dibasic sodium phosphates has long been
established as the method of choice for preparing phosphate buffer. By
mixing the appropriate amounts of monobasic and dibasic sodium phosphate
solutions buffers in the desired pH range can be prepared (see examples on
page 17).
9. Adjust buffer materials to the working pH. Many buffers are supplied as
crystalline acids or bases. The pH of these buffer materials in solution will
not be near the pK a
, and the materials will not exhibit any buffering
capacity until the pH is adjusted. In practice, a buffer material with a pK a
near the desired working pH is selected. If this buffer material is a free acid,
pH is adjusted to desired working pH level by using a base such as sodium
hydroxide, potassium hydroxide, or tetramethyl-ammonium hydroxide.
Alternatively, pH for buffer materials obtained as free bases must be adjusted
by adding a suitable acid.
10. Use buffers without mineral cations when appropriate. Frequently,
buffers without mineral cations are appropriate. Tetramethylammonium
hydroxide fits this criterion. The basicity of this organic quaternary amine is
equivalent to that of sodium or potassium hydroxide. Buffers prepared with
this base can be supplemented at will with various inorganic cations during
the evaluation of mineral ion effects on enzymes or other bioparticulate
activities.
11. Use a graph to calculate buffer composition. Figure 6 shows the
theoretical plot of ∆pH versus [A - ]/[HA] on two-cycle semilog paper. As most
commonly used buffers exhibit only trivial deviations from theoretical value
in the pH range, this plot can be of immense value in calculating the relative
amounts of buffer components required for a particular pH.
For example, suppose one needs 0.1 M MOPS buffer, pH 7.6 at 20°C. At
20°C, the pK a
for MOPS is 7.2. Thus, the working pH is about 0.4 pH units
above the reported pK a
. According to the chart presented, this pH corresponds
to a MOPS sodium/MOPS ratio of 2.5, and 0.1 M solutions of MOPS
and MOPS sodium mixed in this ratio will give the required pH. If any
significant deviations from theoretical values are observed one should check
the proper working conditions and specifications of their pH meter. The
graph can also be used to calculate the amount of acid (or base) required to
adjust a free base buffer material (or free acid buffer material) to the desired
working pH.
17

10
987
6
5
4
3
2
[A - ]/[HA]
1
0.9
0.8
0.7
0.6
0.5
0.4
0.3
0.2
0.1
0.1 -0.8 -0.6 -0.4 -0.2 pK a 0.2 0.4 0.6 0.8 1.0
∆ pH from pK a
Figure 6: Theoretical plot of DpH versus [A - ]/[HA] on two-cycle semilog paper.
Preparation of Some Common Buffers for Use in Biological
Systems
The information provided below is intended only as a general guideline. We
strongly recommend the use of a sensitive pH meter with appropriate temperature
setting for final pH adjustment. Addition of other chemicals, after adjusting
the pH, may change the final pH value to some extent. The buffer concentrations
in the tables below are used only as examples. You may select higher or
lower concentrations depending upon your experimental needs.
1. Hydrochloric Acid-Potassium Chloride Buffer (HCl-KCl); pH Range 1.0
to 2.2
(a) 0.1 M Potassium chloride : 7.45 g/l (M.W.: 74.5)
(b) 0.1 M Hydrochloric acid
Mix 50 ml of potassium chloride and indicated volume of hydrochloric acid.
Mix and adjust the final volume to 100 ml with deionized water. Adjust the final
pH using a sensitive pH meter.
ml of HCl 97 64.5 41.5 26.3 16.6 10.6 6.7
pH 1.0 1.2 1.4 1.6 1.8 2.0 2.2
18

2. Glycine-HCl Buffer; pH range 2.2 to 3.6
(a) 0.1 M Glycine: 7.5 g/l (M.W.: 75.0)
(b) 0.1 M Hydrochloric acid
Mix 50 ml of glycine and indicated volume of hydrochloric acid. Mix and adjust
the final volume to 100 ml with deionized water. Adjust the final pH using a
sensitive pH meter.
ml of HCl 44.0 32.4 24.2 16.8 11.4 8.2 6.4 5.0
pH 2.2 2.4 2.6 2.8 3.0 3.2 3.4 3.6
3. Citrate Buffer; pH range 3.0 to 6.2
(a) 0.1 M Citric acid: 19.21 g/l (M.W.: 192.1)
(b) 0.1 M Sodium citrate dihydrate: 29.4 g/l (M.W.: 294.0)
Mix citric acid and sodium citrate solutions in the proportions indicated and
adjust the final volume to 100 ml with deionized water. Adjust the final pH using
a sensitive pH meter. The use of pentahydrate salt of sodium citrate is not
recommended.
ml of Citric acid 46.5 40.0 35.0 31.5 25.5 20.5 16.0 11.8 7.2
ml of Sodium citrate 3.5 10.0 15.0 18.5 24.5 29.5 34.0 38.2 42.8
pH 3.0 3.4 3.8 4.2 4.6 5.0 5.4 5.8 6.2
4. Acetate Buffer; pH range 3.6 to 5.6
(a) 0.1 M Acetic acid (5.8 ml made to 1000 ml)
(b) 0.1 M Sodium acetate; 8.2 g/l (anhydrous; M.W. 82.0) or 13.6 g/l
(trihydrate; M.W. 136.0)
Mix acetic acid and sodium acetate solutions in the proportions indicated and
adjust the final volume to 100 ml with deionized water. Adjust the final pH using
a sensitive pH meter.
ml of Acetic acid 46.3 41.0 30.5 20.0 14.8 10.5 4.8
ml of Sodium acetate 3.7 9.0 19.5 30.0 35.2 39.5 45.2
pH 3.6 4.0 4.4 4.8 5.0 5.2 5.6
19

5. Citrate-Phosphate Buffer; pH range 2.6 to 7.0
(a) 0.1 M Citric acid; 19.21 g/l (M.W. 192.1)
(b) 0.2 M Dibasic sodium phosphate; 35.6 g/l (dihydrate; M.W. 178.0)
or 53.6 g/l (heptahydrate; M.W. 268.0)
Mix citric acid and sodium phosphate solutions in the proportions indicated and
adjust the final volume to 100 ml with deionized water. Adjust the final pH using
a sensitive pH meter.
ml of Citric acid 44.6 39.8 35.9 32.3 29.4 26.7 24.3 22.2 19.7 16.9 13.6 6.5
ml of Sodium
phosphate
5.4 10.2 14.1 17.7 20.6 23.3 25.7 27.8 30.3 33.1 36.4 43.6
pH 2.6 3.0 3.4 3.8 4.2 4.6 5.0 5.4 5.8 6.2 6.6 7.0
6. Phosphate Buffer; pH range 5.8 to 8.0
(a) 0.1 M Sodium phosphate monobasic; 13.8 g/l (monohydrate, M.W. 138.0)
(b) 0.1 M Sodium phosphate dibasic; 26.8 g/l (heptahydrate, M.W. 268.0)
Mix Sodium phosphate monobasic and dibasic solutions in the proportions
indicated and adjust the final volume to 200 ml with deionized water. Adjust the
final pH using a sensitive pH meter.
ml of Sodium
phosphate, Monobasic
92.0 81.5 73.5 62.5 51.0 39.0 28.0 19.0 13.0 8.5 5.3
ml of Sodium
phosphate, Dibasic
8.0 18.5 26.5 37.5 49.0 61.0 72.0 81.0 87.0 91.5 94.7
pH 5.8 6.2 6.4 6.6 6.8 7.0 7.2 7.4 7.6 7.8 8.0
7. Tris-HCl Buffer, pH range 7.2 to 9.0
(a) 0.1 M Tris(hydroxymethyl)aminomethane; 12.1 g/l (M.W.: 121.0)
(b) 0.1 M Hydrochloric acid
Mix 50 ml of Tris(hydroxymethyl)aminomethane and indicated volume of
hydrochloric acid and adjust the final volume to 200 ml with deionized water.
Adjust the final pH using a sensitive pH meter.
ml of HCl 44.2 41.4 38.4 32.5 21.9 12.2 5.0
pH 7.2 7.4 7.6 7.8 8.2 8.6 9.0
20

8. Glycine-Sodium Hydroxide, pH 8.6 to 10.6
(a) 0.1 M Glycine; 7.5 g/l (M.W.: 75.0)
(b) 0.1 M Sodium hydroxide; 4.0 g/l (M.W.: 40.0)
Mix 50 ml of glycine and indicated volume of sodium hydroxide solutions and
adjust the final volume to 200 ml with deionized water. Adjust the final pH using
a sensitive pH meter.
ml Sodium hydroxide 4.0 8.8 16.8 27.2 32.0 38.6 45.5
pH 8.6 9.0 9.4 9.8 10.0 10.4 10.6
9. Carbonate-Bicarbonate Buffer, pH range 9.2 to 10.6
(a) 0.1 M Sodium carbonate (anhydrous), 10.6 g/l (M.W.: 106.0)
(b) 0.1 M Sodium bicarbonate, 8.4 g/l (M.W.: 84.0)
Mix sodium carbonate and sodium bicarbonate solutions in the proportions
indicated and adjust the final volume to 200 ml with deionized water. Adjust the
final pH using a sensitive pH meter.
ml of Sodium carbonate 4.0 9.5 16.0 22.0 27.5 33.0 38.5 42.5
ml of Sodium bicarbonate 46.0 40.5 34.0 28.0 22.5 17.0 11.5 7.5
pH 9.2 9.4 9.6 9.8 10.0 10.2 10.4 10.6
CALBIOCHEM®
Your Source for High Quality
PROTEIN GRADE®
and
ULTROL® GRADE
Detergents for Over 50 Years.
www.calbiochem.com
21

Commonly Used Buffer Media in Biological Research
Krebs-Henseleit bicarbonate buffer, pH 7.4
119 mM NaCl
4.7 mM KCl
2.5 mM CaCl 2
1.2 mM MgSO 4
1.2 mM KH 2
PO 4
25 mM NaHCO 3
pH 7.4 (at 37°C) when equilibrated with 95% O 2
and 5% CO 2
. Adjust the pH
before use.
Hank’s Biocarbonate Buffer, pH 7.4
137 mM NaCl
5.4 mM KCl
0.25 mM Na 2
HPO 4
0.44 mM KH 2
PO 4
1.3 mM CaCl 2
1.0 mM MgSO 4
4.2 mM NaHCO 3
pH 7.4 (at 37°C) when equilibrated with 95% O 2
and 5% CO 2
. Adjust the pH
before use.
Phosphate Buffered Saline (PBS), pH 7.4
150 mM NaCl
10 mM Potassium Phosphate buffer
(1 liter PBS can be prepared by dissolving 8.7 g NaCl, 1.82 g
K 2
HPO 4
.3H 2
O, and 0.23 g KH 2
PO 4
in 1 liter of distilled water.
Adjust the pH before use).
A variation of PBS can also be prepared as follows:
137 mM NaCl
2.7 mM KCl
10 mM Na 2
HPO 4
1.76 mM KH 2
PO 4
22

Tris Buffered Saline (TBS), pH 7.4
10 mM Tris
150 mM NaCl
(1 liter of TBS can be prepared by dissolving 1.21 g of Tris base and
8.7 g of NaCl in 1 liter of distilled water. Adjust the pH before use.
Note: Tris has a pK a
of 8.3. Hence, the buffering capacity at pH 7.4
is minimal compared to phosphate buffer (pK a
= 7.21).
TBST (Tris Buffered saline and TWEEN®-20)
10 mM Tris-HCl, pH 8.0
150 mM NaCl
0.1% TWEEN®-20
Stripping Buffer for Western Blotting Applications
62.5 mM Tris buffer, pH 6.7 to 6.8
2% Sodium dodecyl sulfate (SDS)
100 mM β-Mercaptoethanol
Cell Lysis Buffer
20 mM Tris-HCl (pH 7.5)
150 mM NaCl
1 mM Sodium EDTA
1 mM EGTA
1% TRITON® X-100
2.5 mM Sodium pyrophosphate
1 mM β-Glycerophosphate
1 mM Sodium orthovanadate
1µg/ml Leupeptin
23

Isoelectric Point of Selected Proteins
Protein Organism/Tissue Isoelectric Point
Acetylcholinesterase Electric eel, Electric organ 4.5
α1-Acid Glycoprotein Human serum 1.8
Acid Protease Penicillium duponti 3.9
Aconitase Porcine heart 8.5
Adenosine deaminase Human erythrocytes 4.7 - 5.1
Adenylate cyclase Mouse brain 5.9 - 6.1
Adenylate kinase Rat liver 7.5 - 8.0
Adenylate Kinase Human erythrocytes 8.5 - 9.0
Albumin Human serum 4.6 - 5.3
Alcohol dehydrogenase Horse liver 8.7 - 9.3
Aldehyde dehydrogenase Rat Liver (cytosol) 8.5
Aldolase Rabbit muscle 8.2 - 8.6
Alkaline phosphatase Bovine intestine 4.4
Alkaline phosphatase Human liver 3.9
cAMP-phosphodiesterase Rat brain 6.1
Amylase Guinea Pig pancreas 8.4
Amylase Human saliva 6.2 - 6.5
Arginase Rat liver 9.4
Arginase Human liver 9.2
ATPase (Na + -K + ) Dog heart 5.1
Carbonic Anhydrase Porcine intestine 7.3
Carboxypeptidase B Human pancreas 6.9
Carnitine acetyltransferase Calf liver 6.0
Catalase Mouse liver (particulate) 6.7
Cathepsin B Human liver 5.1
Cathepsin D Bovine spleen 6.7
Choline acetyltransferase Human brain 7.8
α-Chymotrypsin Bovine pancreas 8.8
Collagenase Clostridium 5.5
C-Reactive protein Human 7.4
DNA polymerase Human lymphocytes 4.7
DNase I Bovine 4.7
Dipeptidase Porcine kidney 4.9
Enolase Rat liver 5.9
Epidermal Growth Factor Mouse submaxillary glands 4.6
Erythropoietin Rabbit plasma 4.8 - 5.0
Ferritin Human liver 5.0 - 5.6
α-Fetoprotein Human serum 4.8
Follicle stimulating hormone Sheep pituitary 4.6
Fructose 1,6-diphosphatase Crab muscle 5.9
Galactokinase Human placenta 5.8
β-Galactosidase Rabbit brain 6.3
Glucose-6-phosphate dehydrogenase Human erythrocytes 5.8 - 7.0
β-Glucuronidase Rat liver microsomes 6.7
24

Isoelectric Point of Selected Proteins, cont.
Protein Organism/Tissue Isoelectric Point
γ-Glutamyl transpeptidase Rat hepatoma 3.0
Glutathione S-transferase Rat liver 6.9, 8.1
D-Glyceraldehyde 3-phosphate dehydrogenase Rabbit muscle 8.3
L-Glycerol-3-phosphate dehydrogenase Rabbit kidney 6.4
Glycogen phosphorylase b Human muscle 6.3
Growth hormone Horse pituitary 7.3
Guanylate kinase Human erythrocytes 5.1
Hemoglobin Rabbit erythrocyte 7.0
Hemoglobin A Human erythrocytes 7.0
Hexokinase Yeast 5.3
Insulin Bovine pancreas 5.7
Lactate dehydrogenase Rabbit muscle 8.5
Leucine aminopeptidase Porcine kidney 4.5
Lipase Human pancreas 4.7
Malate dehydrogenase Rabbit heart (cytosol) 5.1
Malic Enzyme Rabbit heart mitochondria 5.4
Myoglobin Horse muscle 6.8, 7.3
Ornithine decarboxylase Rat liver 4.1
Phosphoenolpyruvate carboxykinase Mouse liver 6.1
Phosphofructokinase Porcine liver 5.0
3-Phosphoglycerate kinase Bovine liver 6.4
Phospholipase A Bee venom 10.5
Phospholipase C C. perfringens 5.3
Phosphorylase kinase Rabbit muscle 5.8
Pepsin Porcine stomach 2.2
Plasmin Human plasma 7.0 - 8.5
Plasminogen Human plasma 6.4 - 8.5
Plasminogen proactivator Human plasma 8.9
Prolactin Human pituitary 6.5
Protein kinase A Bovine brain catalytic subunit 7.8
Prothrombin Human plasma 4.6 - 4.7
Pyruvate kinase Rat liver 5.7
Pyruvate kinase Rat muscle 7.5
Renin Human kidney 5.3
Ribonuclease Bovine pancreas 9.3
RNA polymerase II Human HeLa, KB cells 4.8
Superoxide dismutase Pleurotus olearius 7.0
Thrombin Human plasma 7.1
Transferrin Human plasma 5.9
Trypsin inhibitor Soybean 4.5
Trypsinogen Guinea Porcine pancreas 8.7
Tubulin Porcine brain 5.5
Urease Jack bean 4.9
25

Isoelectric Points of Selected Plasma/Serum Proteins
Protein M.W. Species Isoelectric Point
α 1
-Acid Glycoprotein 44,000 Human 2.7
Albumin 66,000 Human 5.2
α 1
-Antitrypsin 51,000 Human 4.2 - 4.7
Ceruloplasmin 135,000 Human 4.4
Cholinesterase 320,000 Human 4.0
Conalbumin — Human 5.9
C-Reactive Protein 110,000 Human 4.8
Erythropoietin — Rabbit 4.8 - 5.0
α-Fetoprotein 70,000 Human 4.8
Fibrinogen 340,000 Human 5.5
IgG 150,000 Human 5.8 - 7.3
IgD 172,000 Human 4.7 - 6.1
β-Lactoglobulin 44,000 Bovine 5.2
α 2
-Macroglobulin 725,000 Human 5.4
β 2
-Macroglobulin 11,800 Human 5.8
Plasmin — Human 7.0 - 8.5
Prealbumin 50,000 - 60,000 Human 4.7
Prothrombin — Bovine 4.6 - 4.9
Thrombin 37,000 Human 7.1
Thyroxine Binding Protein 63,000 Human 4.2 - 5.2
Transferrin 79,600 Human 5.9
Approximate pH and Bicarbonate Concentration in
Extracellular Fluids
Fluid pH meq HCO3¯/liter
Plasma 7.35 - 7.45 28
Cerebrospinal Fluid 7.4 25
Saliva 6.4 - 7.4 10 - 20
Gastric Secretions 1.0 - 2.0 0
Tears 7.0 - 7.4 5 - 25
Aqueous Humor 7.4 28
Pancreatic Juice 7.0 - 8.0 80
Sweat 4.5 - 7.5 0 - 10
26

Ionization Constants K and pK a
for Selected Acids and Bases
in Water
Acids and Bases Ionization Constant (K) pK a
Acetic Acid 1.75 x 10 -5 4.76
Citric Acid 7.4 x 10 -4 3.13
1.7 x 10 -5 4.77
4.0 x 10 -7 6.40
Formic Acid 1.76 x 10 -4 3.75
Glycine 4.5 x 10 -3 2.35
1.7 x 10 -10 9.77
Imidazole 1.01 x 10 -7 6.95
Phosphoric Acid 7.5 x 10 -3 2.12
6.2 x10 -8 7.21
4.8 x 10 -13 12.32
Pyruvic Acid 3.23 x 10 -3 2.49
Tris(hydroxymethyl)aminomethane 8.32 x 10 -9 8.08
Physical Properties of Some Commonly Used Acids
Molecular Specific % Weight/ Approx.
ml required to
Acid
Weight Gravity Weight Normality
make 1 liter
of 1 N solution
Acetic Acid 60.05 1.06 99.50 17.6 57
Hydrochloric Acid 36.46 1.19 37 12.1 83
Nitric Acid 63.02 1.42 70 15.7 64
Perchloric Acid (72%) 100.46 1.68 72 11.9 84
Phosphoric Acid 98.00 1.70 85 44.1 23
Sulfuric Acid 98.08 1.84 96 36.0 28
27

Some Useful Tips for Calculation of Concentrations and
Spectrophotometric Measurements
As per Beer’s law
A = abc
Where A = absorbance
a = proportionality constant defined as absorptivity
b = light path in cm
c = concentration of the absorbing compound
When b is 1 cm and c is moles/liter, the symbol a is substituted by a symbol e
(epsilon).
e is a constant for a given compound at a given wavelength under prescribed
conditions of solvent, temperature, pH and is called as molar absorptivity. e is
used to characterize compounds and establish their purity.
Example:
Bilirubin dissolved in chloroform at 25°C should have a molar absorptivity (e) of
60,700.
Molecular weight of bilirubin is 584.
Hence 5 mg/liter (0.005 g/l) read in 1 cm cuvette should have an absorbance of
A = (60,700)(1)(0.005/584) = 0.52 {A = abc}
Conversely, a solution of this concentration showing absorbance of 0.49 should
have a purity of 94% (0.49/0.52).
In most biochemical and toxicological work, it is customary to list constants
based on concentrations in g/dl rather than mol/liter. This is also common when
molecular weight of the substance is not precisely known.
Here for b = 1 cm; and c = 1 g/dl (1%), A can be written as A
1%
1cm
This constant is known as absorption coefficient.
The direct proportionality between absorbance and concentration must be
established experimentally for a given instrument under specified conditions.
28

Frequently there is a linear relationship up to a certain concentration. Within
these limitations, a calibration constant (K) may be derived as follows:
A = abc.
Therefore,
c = A/ab = A x 1/ab.
The absorptivity (a) and light path (b) remain constant in a given method of
analysis. Hence, 1/ab can be replaced by a constant (K).
Then,
c = A x K; where K = c/A. The value of the constant K is obtained by measuring
the absorbance (A) of a standard of known concentration (c).
29

CALBIOCHEM® Buffers
We offer an extensive line of buffer materials that meet the highest standards of
quality. We are continuing to broaden our line of ULTROL® Grade Buffer
materials, which are of superior quality and are manufactured to meet stringent
specifications. In addition, whenever possible, they are screened for uniform
particle size, giving uniform solubility characteristics.
ADA, Sodium Salt
(N-2-Acetamido-2-iminodiacetic Acid, Na)
M.W. 212.2
100 g
Cat. No. 114801
2-Amino-2-methyl-1,3-propanediol
M.W. 105.1
50 g
Cat. No. 164548
500 g
BES, Free Acid, ULTROL® Grade
[N,N-bis-(2-Hydroxyethyl)-2-aminoethanesulfonic
Acid]
M.W. 213.3
25 g
Cat. No. 391334
Bicine, ULTROL® Grade
[N,N-bis-(2-Hydroxyethyl)glycine]
M.W. 163.2
Cat. No. 391336
100 g
1 kg
BIS-Tris, ULTROL® Grade
{bis(2-Hydroxyethyl)imino]-tris(hydroxymethyl)methane}
M.W. 209.2
100 g
Cat. No. 391335
1 kg
BIS-Tris Propane, ULTROL® Grade
{1,3-bis[tris(Hydroxymethyl)methylamino]-
propane}
M.W. 282.4
100 g
Cat. No. 394111
1 kg
Boric Acid, Molecular Biology Grade
M.W. 61.8
500 g
Cat. No. 203667
1 kg
5 kg
Cacodylic Acid, Sodium Salt
(Sodium Dimethyl Arsenate)
M.W. 160.0
Cat. No. 205541
100 g
CAPS, ULTROL® Grade
[3-(Cyclohexylamino)propanesulfonic Acid]
M.W. 221.3
100 g
Cat. No. 239782
1 kg
CHES, ULTROL® Grade
[2-(N-Cyclohexylamino)ethanesulfonic Acid]
M.W. 207.3
100 g
Cat. No. 239779
Citric Acid, Monohydrate, Molecular Biology
Grade
M.W. 210.1
100 g
Cat. No. 231211
1 kg
Glycine, Free Base
M.W. 75.1
Cat. No. 3570
Glycine, Molecular Biology Grade
M.W. 75.1
Cat. No. 357002
Glycylglycine, Free Base
M.W. 132.1
Cat. No. 3630
500 g
100 g
1 kg
25 g
100 g
HEPES, Free Acid, Molecular Biology Grade
(N-2-Hydroxyethylpiperazine-N′-2-
ethanesulfonic Acid)
M.W. 238.3
25 g
Cat. No. 391340
250 g
HEPES, Free Acid, ULTROL® Grade
(N-2-Hydroxyethylpiperazine-N′-2-
ethanesulfonic Acid)
M.W. 238.3
Cat. No. 391338
25 g
100 g
500 g
1 kg
5 kg
30

HEPES, Free Acid, ULTROL® Grade, 1 M
Solution
M.W. 238.3
100 ml
Cat. No. 375368
500 ml
HEPES, Sodium Salt, ULTROL® Grade
(N-2-Hydroxyethylpiperazine-N′-2-
ethanesulfonic Acid, Na)
M.W. 260.3
Cat. No. 391333
100 g
500 g
1 kg
HEPPS, ULTROL® Grade
(EPPS; N-2-Hydroxyethylpiperazine-N′-3-
propane sulfonic acid)
M.W. 252.3
100 g
Cat. No. 391339
Imidazole, ULTROL® Grade
(1,3-Diaza-2,4-cyclopentadiene )
M.W. 68.1
Cat. No. 4015
25 g
100 g
MES, Free Acid, ULTROL® Grade
[2-(N-Morpholino)ethanesulfonic Acid]
M.W. 195.2
100 g
Cat. No. 475893
500 g
1 kg
MES, Sodium Salt, ULTROL® Grade
[2-(N-Morpholino)ethanesulfonic Acid, Na]
M.W. 217.2
100 g
Cat. No. 475894
1 kg
MOPS, Free Acid, ULTROL® Grade
[3-(N-Morpholino)propanesulfonic Acid]
M.W. 209.3
100 g
Cat. No. 475898
500 g
1 kg
MOPS, Sodium, ULTROL® Grade
[3-(N-Morpholino)propanesulfonic Acid, Na]
M.W. 231.2
100 g
Cat. No. 475899
1 kg
MOPS/EDTA Buffer, 10X Liquid Concentrate,
Molecular Biology Grade
M.W. 209.3
100 ml
Cat. No. 475916
PBS-TWEEN® Tablets
(Phosphate Buffered Saline-TWEEN® 20 Tablets)
Cat. No. 524653
1 each
PIPES, Free Acid, Molecular Biology Grade
[Piperazine-N,N′-bis(2-ethanesulfonic Acid)]
M.W. 302.4
25 g
Cat. No. 528133
250 g
PIPES, Free Acid, ULTROL® Grade
[Piperazine-N,N′-bis(2-ethanesulfonic Acid)]
M.W. 302.4
100 g
Cat. No. 528131
1 kg
PIPES, Sesquisodium Salt, ULTROL® Grade
[Piperazine-N,N′-bis(2-ethanesulfonic Acid),
1.5Na]
M.W. 335.3
100 g
Cat. No. 528132
1 kg
PIPPS
[Piperazine-N,N′-bis(3-propanesulfonic Acid)]
M.W. 330.4
10 g
Cat. No. 528315
Potassium Phosphate, Dibasic, Trihydrate,
Molecular Biology Grade
M.W. 228.2
250 g
Cat. No. 529567
1 kg
Potassium Phosphate, Monobasic
M.W. 136.1
Cat. No. 529565
100 g
500 g
Potassium Phosphate, Monobasic, Molecular
Biology Grade
M.W. 136.1
250 g
Cat. No. 529568
1 kg
Sodium Citrate, Dihydrate
(Citric Acid, 3Na)
M.W. 294.1
Cat. No. 567444
Sodium Citrate, Dihydrate, Molecular
Biology Grade
M.W. 294.1
Cat. No. 567446
1 kg
100 g
1 kg
5 kg
PBS Tablets
(Phosphate Buffered Saline Tablets)
Cat. No. 524650
1 each
Sodium Phosphate, Dibasic
M.W. 142.0
Cat. No. 567550
500 g
1 kg
31

Sodium Phosphate, Dibasic, Molecular
Biology Grade
M.W. 142.0
250 g
Cat. No. 567547
1 kg
Sodium Phosphate, Monobasic
M.W. 120.0
Cat. No. 567545
250 g
500 g
1 kg
Sodium Phosphate, Monobasic, Monohydrate,
Molecular Biology Grade
M.W. 138.0
250 g
Cat. No. 567549
1 kg
Triethanolamine, Hydrochloride*
M.W. 185.7
Cat. No. 641752
1 kg
Triethylammonium Acetate, 1 M Solution
M.W. 161.2
1 liter
Cat. No. 625718
Tris Base, Molecular Biology Grade
[tris(Hydroxymethyl)aminomethane]
M.W. 121.1
Cat. No. 648310
100 g
500 g
1 kg
2.5 kg
SSC Buffer, 20X Powder Pack, ULTROL®
Grade
Cat. No. 567780
2 pack
SSPE Buffer, 20X Powder Pack, ULTROL®
Grade
Cat. No. 567784
2 pack
Tris Base, ULTROL® Grade
[tris(Hydroxymethyl)aminomethane]
M.W. 121.1
Cat. No. 648311
100 g
500 g
1 kg
5 kg
10 kg
TAPS, ULTROL® Grade
(3-{[tris(Hydroxymethyl)methyl]amino}-
propanesulfonic Acid)
M.W. 243.2
100 g
Cat. No. 394675
1 kg
TBE Buffer, 10X Powder Pack, ULTROL®
Grade
(10X Tris-Borate-EDTA Buffer)
Cat. No. 574796
2 pack
TES, Free Acid, ULTROL® Grade
(2-{[tris(Hydroxymethyl)methyl]amino}-
ethanesulfonic Acid)
M.W. 229.3
100 g
Cat. No. 39465
1 kg
TES, Sodium Salt, ULTROL® Grade
M.W. 251.2
Cat. No. 394651
100 g
Tricine, ULTROL® Grade
{N-[tris(Hydroxymethyl)methyl]glycine}
M.W. 179.2
100 g
Cat. No. 39468
1 kg
Tris Buffer, 1.0 M, pH 8.0, Molecular Biology
Grade
M.W. 121.1
100 ml
Cat. No. 648314
Tris Buffer, 100 mM, pH 7.4, Molecular
Biology Grade
M.W. 121.1
100 ml
Cat. No. 648315
Tris, Hydrochloride, Molecular Biology
Grade
[tris(Hydroxymethyl)aminomethane, HCl]
M.W. 157.6
100 ml
Cat. No. 648317
1 kg
Tris, Hydrochloride, ULTROL® Grade
[tris(Hydroxymethyl)aminomethane, HCl]
M.W. 157.6
250 g
Cat. No. 648313
500 g
1 kg
* Not for international sales outside the US.
Buy in Bulk and $ave!
To request a bulk quotation,
call our bulk department at
800-854-2855 or your local sales office.
32